Chemistry - Chapter Summaries - John Green and Sadru Damji - Third Edition - IBID 2008 PDF

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CHAPTER SUMMARIES

CHAPTER 1 QUANTITATIVE CHEMISTRY (IB TOPIC 1) SUMMARY Introduction •

1 dm3 = 1 litre = 1 x 10-3m3 = 1 x 103 cm3 = 1000 ml

•

Amount of substance, n, is measured in moles (mol).

•

1 mol of a chemical species contains the same number of particles as there are atoms in exactly 12 g of C-12 ( 126 C ) isotope.

•

1 mol of any substance contains 6.02 × 1023 particles; 6.02 × 1023 mol−1 is called Avogadro’s Constant (L or NA)

•

Ar, the relative atomic mass of an element is the ratio of the mass of an atom of the element to the mass of one atom of C-12; Ar has no units.

•

Mr, the relative molecular mass also has no units.

•

M, the molar mass is the mass of one mole of any chemical species and has the units g mol−1.

The amount of substance, n in moles =

mass ( g ) ; molar mass ( g mol −1 )

V ( dm3 ) . n= Vm ( dm 3 mol −1 ) ⇔ Number of Particles

Avogadro Constant (6.02×1023 mol−1)

Amount

⇔

Mass

in moles

Molar mass

in grams

•

Use Avogadro Constant to convert between amount and number of particles.

•

Use molar mass of substance to convert between amount in moles and mass in grams.

•

Empirical formula gives the simplest whole number ratio of atoms in a compound.

•

Molecular formula gives the actual number of atoms of each element in the molecule of a compound.

If a molecular formula is given, percentage composition can be calculated.

If the percentage composition is given: •

Consider 100 g of sample, the % of each element becomes its mass.

•

Convert the mass of each element to its amount in moles.

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Determine simplest whole number ratio – this is its empirical formula.

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Molecular formula is a whole-number multiple of the empirical formula.

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Substances react by amounts based on a balanced chemical equation.

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CHAPTER 1 QUANTITATIVE CHEMISTRY (IB TOPIC 1) SUMMARY If masses are given, convert these to amounts of substances. •

In stoichiometric problems, apply mole ratios as specified in the balanced chemical equation.

•

The % yield indicates how efficient a reaction is (given by:

experimental yield × 100% ). theoretical yield •

STP for gases is standard temperature (0°C or 273 K) and pressure (1 atmosphere or 101.325 kPa).

•

Molar volume, Vm, of any gas at STP = 22.4 dm3.

In a balanced chemical equation, coefficients stand for the amount of substance. For gases, these also refer to volumes of gases. Concentration, c =

amount n (mol) n = mol dm −3 . 3 Vol solution (dm ) V

n = cV; volume must be in dm3. On dilution, the amount of solute does not change but the volume increases and the concentration decreases Amount n =

m (g) . M (g mol −1 )

For a reaction aA + bB → products, where a and b are coefficients, then 1 1 a n A = b nB. Percentage yield =

experimental yield × 100%. theoretical yield

(There are many worked examples given in the chapter)

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CHAPTER 2 ATOMIC STRUCTURE (IB TOPICS 2 AND 12) SUMMARY Introduction •

Relative masses: p = 1, n = 1, e = 1/1840; charges: p = +1, n = 0, e– = –1.

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Protons and neutrons are present in the nucleus of an atom, electrons are in orbits or shells around the nucleus.

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Atomic number, Z = number of protons; the fundamental characteristic of an element.

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Mass number, A = number of (protons + neutrons).

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Isotopes: same atomic number, different mass number OR same number of protons, different number of neutrons OR atoms of the same element with different masses.

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Isotopes differ in physical properties that depend on mass such as density, rate of diffusion etc. Chemical properties are the same because of the same electronic configuration or arrangement.

•

Atomic mass of an element is the average of the atomic masses of its isotopes; depends on isotopes relative abundance; leads to non-integer atomic masses.

Mass Spectrometer •

Stages of Operation: Vaporization of sample, ionization to produce M+ ions, acceleration of ions by electric field, deflection of ions by magnetic field, vacuum, detection of ions.

•

Degree of deflection:

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Lower the mass, higher the deflection.

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Higher the charge, higher the deflection.

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Deflection reflects mass/charge ratio; for charge of +1, deflection depends on mass.

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For an element, the mass spectrum gives two important pieces of information: the number of isotopes, and the abundance of each isotope; thus the relative average atomic mass, Ar can be calculated.

•

For a molecule, the highest peak represents the molecular (parent) ion and its mass gives the relative molecular mass, Mr of the compound (and the fragmentation pattern can help determine its structure). o

A continuous spectrum contains light of all wavelengths in the visible range.

o

A line spectrum consists of a few lines of different wavelengths.

o

When electrons are excited, they jump to higher energy levels.

o

Electrons fall back to lower energy levels, and the energy equivalent to the difference in energy level is emitted in the form of photons.

o

Energy levels come together in terms of energy the farther away they are from the nucleus; this explains the convergence of lines in a line spectrum.

o

The maximum number of electrons in a main energy level n is 2n2: 1st energy level, n = 1; maximum 2 e–; n = 2, maximum 8 e–; n = 3, maximum 18 e–.

The electron arrangement (or configuration) indicates the number of electrons and their energy distribution. This determines an element’s physical and chemical properties.

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CHAPTER 2 ATOMIC STRUCTURE (IB TOPICS 2 AND 12) SUMMARY •

Main (or principal) energy levels, sub-levels and orbitals: The main energy levels, n are assigned whole number integers, n = 1, 2, 3, 4… . n = 1 represents the lowest energy level. Each main energy level contains n sub-levels and a total of n2 orbitals. o

s, p, d, f etc. is the common notation for sub-levels and orbitals within sub-levels. An orbital is an area of space around the nucleus in which an electron moves.

o

Orbitals have characteristic shapes. There is one s orbital which is spherical in shape, three p orbitals which are dumbbell shaped, called px, py pz, and arranged in the x, y, and z directions respectively, five d orbitals and seven f orbitals (both with complex shapes). The relative energies of s, p, d, and f orbitals with in a sub-level are: s < p < d < f.

o

Each orbital can have a maximum of 2 electrons. n = 1 has one sub-level which is called an s sub-level and which contains one s orbital. n = 2 has two sub-levels: 2s and 2p; n = 3 has 3 sub-levels: 3s, 3p and 3d; n = 4 has 4 sub-levels:4s, 4p, 4d and 4f, etc.

•

The Aufbau (‘building-up’) Principle: Electrons are placed in orbitals in order of increasing energy, starting with the lowest energy level, and in general, filling each sublevel completely before beginning the next. This is due to the fact that systems in nature prefer minimum energy in order to achieve maximum stability.

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Hund’s Rule: Occupation of sub-levels takes place singly as far as possible before pairing starts.

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Pauli exclusion principle: No two electrons in an atom can be in exactly the same state; no two electrons in a given atom can have the same four quantum numbers (that is, these can not be in the same place at the same time)

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nlx notation is used to describe the electron configuration of an element: n is the main energy level, l the sub-level, and x is the number of electrons in the sub-level.

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The ionisation energy of an atom is the minimum amount of energy required to remove a mole of electrons from a mole of gaseous atoms to form a mole of gaseous ions.

(N.B. Shading indicates Topic 12 (AHL) material.)

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CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY Basic concepts •

Periodic Table is arranged according to increasing atomic number and consists of horizontal rows called periods, and vertical columns called groups (or families). The arrangement is such that elements with similar chemical properties fall directly beneath each other in the same group.

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A period is a series of elements arranged according to increasing atomic number, which begins with the first element having one electron in a new main energy level.

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A group is a vertical column consisting of elements with the same electron arrangement in their outer energy levels, which gives the group similar chemical properties.

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Modern Periodic Law states that chemical and physical properties of elements vary periodically if elements are arranged in order of increasing atomic numbers.

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Periodicity is the regular repeating of properties according to arrangement of elements in the periodic table (i.e., after regular intervals) such as atomic radius, ionisation energy, etc. arising from the systematic filling of successive energy level.

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Valence electrons are electrons in the outermost energy level (the highest energy level) of an atom and are usually the electrons that take part in a chemical reaction.

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Properties of elements are chiefly due to the number and arrangement of electrons in the outer energy level of atoms.

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The number of valence electrons is the same for a group, but increases across a period.

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The transition metals are the large d-block elements in the middle of the Periodic Table from Sc to Zn etc. (includes three transition series).

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Ionization Energy (IE) is the minimum energy required to form a mole of 1+ ions (by removing an electron from each atom) in the gaseous state: M (g) Æ M+ (g) + e– (units: kJ mol–1)

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Electronegativity is a measure of how strongly the atom attracts the electrons in a covalent bond.

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CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY

Alkali metals  are extremely reactive, electropositive metals. They react vigorously with O2 in the air and oxidize fast to form ionic oxides.  contain one valence electron that is very easily lost due to low ionization energy; thus they are very good reducing agents: Na (s) → Na+ (aq) + e–.  In compounds, these exist as +1 cations with noble gas electron configuration.  Since a valence e– further away from the nucleus down a group is more easily removed, reactivity increases from Li to Cs.  Reaction with Water: Reactivity increases down the group: 2 Li (s) + 2 H2O (l) Æ 2 LiOH (aq) + H2 (g) + heat •

Reaction with Halogens: form ionic salts (e.g. Na+Cl–); the reaction is highly exothermic. The alkali halides are ionic, neutral, water-soluble, white crystalline compounds. 2 Na (s) + Cl2 (g) Æ 2 NaCl (s) + heat

Halogens 

Are non-metallic; metallic character increases down the group (I2 is a shiny solid).



Higher ionization energies indicate little tendency to lose electrons.



Have 7 valence electrons and achieve noble gas configuration by gaining an electron to form an anion (e.g. Cl–), or by sharing an e– pair with another atom (e.g., Cl2).



Are diatomic (F2, Cl2, Br2, I2), non-polar molecules. They are simple molecular substances with only weak van der Waal’s between molecules (Bonding, Ch. 4).



Higher electronegativities mean that halogens have a tendency to accept an electron, and act as oxidizing agents: X2 +2 e– Æ 2 X–.



Oxidizing strength decreases down the group, since the atom gets larger and attraction for the electrons decreases.



Reactions of halogens X2 with halide ions X–: Halogen Displacement Reactions: Weaker halogens are displaced from their salts by more powerful oxidizing agents. Thus Cl2 can displace Br2 from Br– and I2 from I– (as it is a stronger oxidizing agent than both Br and I), but it cannot displace F2 from F–.



Identification of Halide Ions – Reaction with silver ions, Ag+: Silver halides, with different colored precipiates can be used to identify a halide ion: 1. AgNO3 (aq) + NaF (aq) → No precipitate formed. 2. AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq), a white precipitate. 3. Ag+ (aq) + Br– (aq) → AgBr (s), a cream precipitate. 4. Ag+ (aq) + I– (aq) → AgI (s), a yellow precipitate.

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CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY

Period 3 elements  Ionization energy and electronegativity increase across the period; thus elements do not lose electrons as easily and metallic character decreases across the period. Bonding changes from metallic to covalent. • Na, Mg and Al are metallic, good conductors of heat and electricity, with low to medium melting points, and lower electronegativities. • Si is a semi-conductor / metalloid. It forms a network covalent solid of very high melting point. • P, S and Cl are elements of higher electronegativity and are non-metallic. These form simple molecular substances (see bonding) with lower boiling points; the bonding between atoms is covalent and the bonding between molecules is weak van der Waal’s forces.

Period 3 oxides and chlorides Acid/base properties of oxides

Acid/base properties of chlorides

Ionic solids with high melting points; bonding between active metals of low IE and active nonmetals of high electronegativity

Ions present, e.g. Oxides form basic solutions Na+ + Cl− Mg2+ + O2−; conduct electricity in the molten state (and in aqueous sol)

Na2O + H2O Æ 2NaOH;

Chlorides (NaCl, MgCl2) form neutral solutions. MgCl2 is actually very slightly acidic.

Al2O3 is ionic; SiO2 is a giant covalent structure (with very high melting point).

No mobile ions present; nonconductors

Oxides are acidic except Al2O3 which is amphoteric (can act as an acid or a base)

Period 3 oxides and chlorides

Bonding in oxides and chlorides

s block elements: Na and Mg

p block elements Al, Si; P, S, Cl

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Others: covalent bonding; form simple molecular substances

Electrical conductivityof molten oxides and chlorides

MgO + H2O Æ Mg(OH)2

Chlorides are acidic; undergo hydrolysis reaction to produce HCl vapor

pH of chlorides in water

NaCl : pH =7 MgCl2: pH ∼ 7

Al2Cl6: pH ∼ 3 SiCl4, PCl5, S2Cl2 : pH ∼ 2

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CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY d-block elements •

d-block elements with characteristic properties form at least one ion with a partially filled d sub-level (with 1 to 9 electrons); Sc and Zn are not typical of d-block elements.

•

Characteristic properties are: presence of variable oxidation states, formation of complex ions, colored complexes, and catalytic activity.

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4s is lower in energy (but further from the nucleus) than 3d sub-level; 4s electrons are lost before 3d; both 4s and 3d can behave as valence electrons because these are close in energy.

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Multiple oxidation states are due to 4s and 3d being close in energy. All d-block elements (except Sc) show an oxidation state of +2. All d-block elements except Zn show an oxidation state of +3.

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Complexes are formed when a central metal ion is bonded to ligands; ligands contain (at least one) lone e− pair that form coordinate bonds with the central transition metal ion.

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Color is due to d-d electron transitions between split d orbitals.

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Transition metals behave as surface catalysts and their compounds more often as intermediate catalysts, which are important for industrial and biological reactions.

(N.B. Shading indicates Topic 13 (AHL) material.)

© IBID Press 2007

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CHAPTER 4 BONDING (IB TOPICS 4 AND 14) SUMMARY A chemical bond is the interaction between atoms within a molecule, between molecules or between ions of opposite charges. Bond formation is an exothermic process; it gives out energy and leads to a more energetically stable state. Bonding in liquids and solids 1. Covalent bonding results from electron sharing between non-metals or a non-metal and a metal of higher electronegativity. The electron pair is attracted by both nuclei leading to a bond that is directional in nature. (i) van der Waals’ Forces: For a non-polar molecule, only weak, temporary, instantaneous dipole-dipole interaction called van der Waals’ forces exist between molecules. These molecules are low melting point solids and low boiling point liquids and gases. Larger the molecule, stronger the van der Waals’ forces, higher its boiling and melting points. (ii) Dipole-dipole Interaction: In a polar molecule, besides weak van der Waals’ forces, the molecules experience stronger permanent dipole-dipole interaction. (iii) Hydrogen Bonding: Elements of high electronegativity (F, O, N), bonded to a (tiny) hydrogen atom give rise to a special case of dipole-dipole interaction called H-bonding. This is important in determining solubility, melting and boiling points, and stability of crystal structures. Hydrogen bonding also plays an important role in biological systems. Hydrogen bonded molecules experience stronger hydrogen bonding in addition to van der Waals’ forces and dipole-dipole interaction. Strength of bonding: van der Waals’ < dipole-dipole < H-bond 1, reaction goes almost to completion. • •

When KC < 1, concentration of reactants generally exceeds that of products at equilibrium. When KC 7; use phenolphthalein (pH range 8 – 10) (c) Strong acid + weak base: solution at equivalence point, pH < 7; use methyl orange (pH range: 3.1 – 4.4) (d) Weak acid + weak base: solution at equivalence point, pH about 7; gradual change in pH; no suitable color indicator; record change in pH using a pH meter and obtain point of inflection for equivalence point.

Methods of Determining Ka of a Weak Acid or a Weak Base 1. Determine pH of a solution of known concentration of the weak acid. Knowing pH, calculate [H+] = x. HX (aq) ' H+ (aq) + X– (aq) [initial]: 0.10 — — [eq.]: Ka =

0.10 – x

x

x

Assume x is negligible; i.e., x C–Br > C–Cl > C–F, because as the size increases down the halogen group, the bonding electrons in C–I are the furthest apart. It is therefore the longest and weakest bond whereas C–F would be the shortest and strongest bond. Thus the halide leaves the fastest in the order I– > Br– > Cl– > F– and the rate ...